Basic Chemistry

Water
is an odorless, tasteless, transparent liquid that is colorless in small amounts but exhibits a bluish tinge in large quantities. It is the most familiar and abundant liquid on earth. In solid form ( ice ) and liquid form it covers about 70% of the earth's surface. It is present in varying amounts in the atmosphere. Most of the living tissue of a human being is made up of water; it constitutes about 92% of blood plasma, about 80% of muscle tissue, about 60% of red blood cells, and over half of most other tissues. It is also an important component of the tissues of most other living things.

The word water comes from the Old English waeter; GermanWasser, originally from hypothetical Proto-Indo-European wod-or.

Chemical and Physical Properties

Chemically, water is a compound of hydrogen and oxygen, having the formula H2O. It is chemically active, reacting with certain metals and metal oxides to form bases, and with certain oxides of nonmetals to form acids. It reacts with certain organic compounds to form a variety of products, e.g., alcohols from alkenes. Because water is a polar compound, it is a good solvent. Although completely pure water is a poor conductor of electricity, it is a much better conductor than most other pure liquids because of its self-ionization, i.e., the ability of two water molecules to react to form a hydroxide ion, OH-, and a hydronium ion, H3O+. Its polarity and ionization are both due to the high dielectric constant of water.

Water has interesting thermal properties. When heated from 0°C, its melting point, to 4°C, it contracts and becomes more dense; most other substances expand and become less dense when heated. Conversely, when water is cooled in this temperature range, it expands. It expands greatly as it freezes; as a consequence, ice is less dense than water and floats on it. Because of hydrogen bonding between water molecules, the latent heats of fusion and of evaporation and the heat capacity of water are all unusually high. For these reasons, water serves both as a heat-transfer medium ( e.g., ice for cooling and steam for heating ) and as a temperature regulator ( the water in lakes and oceans helps regulate the climate. )

Structure of the Water Molecule

Many of the physical and chemical properties of water are due to its structure. The atoms in the water molecule are arranged with the two H-O bonds at an angle of about 105 degrees rather than on directly opposite sides of the oxygen atom. The asymmetrical shape of the molecule arises from a tendency of the four electron pairs in the valence shell of oxygen to arrange themselves symmetrically at the vertices of a tetrahedron around the oxygen nucleus. The two pairs associated with covalent bonds holding the hydrogen atoms are drawn together slightly, resulting in the angle of 105 degrees between these bonds. This arrangement results in a polar molecule, since there is a net negative charge toward the oxygen end ( the apex ) of the V-shaped molecule and a net positive charge at the hydrogen end. The electric dipole gives rise to attractions between neighboring opposite ends of water molecules, with each oxygen being able to attract two nearby hydrogen atoms of two other water molecules. Such hydrogen bonding, as it is called, has also been observed in other hydrogen compounds. Although considerably weaker than the covalent bonds holding the water molecule together, hydrogen bonding is strong enough to keep water liquid at ordinary temperatures; its low molecular weight would normally tend to make it a gas at such temperatures.

Various other properties of water, such as its high specific heat, are due to these hydrogen bonds. As the temperature of water is lowered, clusters of molecules form through hydrogen bonding, with each molecule being linked to others by up to four hydrogen bonds, each oxygen atom tending to surround itself with four hydrogen atoms in a tetrahedral arrangement. Hexagonal rings of oxygen atoms are formed in this way, with alternate atoms in either a higher or lower plane than their neighbors to create a kinked three-dimensional structure.

Liquid Water

According to present theories, water in the liquid form contains three different molecule populations. At the highest temperatures single molecules are the rule, with little hydrogen bonding because of the high thermal energy of the molecules. In the middle range of temperatures there is more hydrogen bonding, and clusters of molecules are formed. At lower temperatures aggregates of clusters also form, these aggregates being the most common arrangement below about 15°C.

On the basis of these three population types and the transitions between them, many aspects of the anomalous behavior of water can be explained. For example, the tendency of water to freeze faster if it has been cooled rapidly from a relatively warm temperature than if it has been cooled at the same rate from a lower temperature is explained in terms of the greater number of irregularly shaped cluster aggregates in the cooler water that must find a suitable means of fitting together with a neighboring aggregate.

Ice

In ice, each molecule forms the maximum number of hydrogen bonds, resulting in crystals composed of open, hexagonal columns. Because these crystals have a number of open regions and pockets, normal ice is less dense than water. However, other forms of ice also exist at conditions of higher pressure, each of these different forms ( designated ice II, ice III, etc. ) having greater density and other distinct physical properties that differ from those of normal ice, or ice I. As many as eight different forms of ice have been distinguished in this manner. * the higher pressures creating such forms cause rearrangements of the hexagonal columns in ice, although the basic kinked hexagonal ring is common to all forms.

When ice melts, it is thought that the fragments of these structures fill many of the gaps that existed in the crystal lattice, making water denser than ice. This tendency is the dominant one between 0°C and 4°C, at which temperature water reaches its maximum density. Above this temperature, expansion due to the increased thermal energy of the molecules is the dominant factor, with a consequent decrease in density.

* Editor's note: I recall once reading a sci-fi short story in which scientists created "Ice IX" - the ninth form of ice. This artificial ice had two interesting properties: first, it was its own catalyst - self-reproducing, and second, it could form at temperatures well above freezing. Of course, when it got out of the lab, the whole earth "froze".

Density

Pure liquid water has a density of 62.4 pounds per cubic foot ( or 8.341pounds per gallon ) at a reference temperature of 39°F. Over a biologically relevant range of temperatures, the density of pure liquid water ranges from 62.42 lbs/cu.ft at 32°F to about 62.0 lbs/cu.ft at 85°F. The density of seawater is somewhat higher owing to the dissolved salts, ranging from 64.1 to 64.9, with a standard value of 64.6 at 39°F. The density of seawater is highly dependant on the salinity.

The density of ice is considerably less at 56.2 pounds per cubic foot; water being one of the few liquids that shows a decrease in density upon solidifying. Thus, ice floats. While this was unfortunate for the Titanic, it is otherwise a very lucky thing for life on earth. If ice did not float, all bodies of water would freeze solid from the bottom up, and life as we know it would have a very difficult time surviving! As it is, a winter surface layer of ice and snow actually insulates the liquid water below and protects the life in it.

Viscosity

Temperature
( oF )

Absolute
Dynamic
Viscosity
( cp )

32

1.794

40

1.546

50

1.310

60

1.129

70

0.982

80

0.862

90

0.764

Viscosity is the measure of a liquid's resistance to flow, or its opposition to the movement of an object through it. High viscosity liquids are "thick", like syrups and motor oil. Water is a fairly low viscosity liquid. The viscosity of water varies with temperature and density, although the temperature variation is far greater than the density variation.

The viscosity of fresh water over a range of biologically relevant temperatures is shown at right. Saltwater is slightly higher. Without going into too much detail, it is apparent that cold water is approximately twice as viscous as warm water, or that drag in cold water is much greater than in warm water.

This in turn explains why warm water fishes can have such outlandish shapes, while cold water fishes are all much more streamlined and boring. This may also explain why you can swim faster in the Bahamas than in New Jersey, although I think a drysuit and double tanks might also have something to do with it.

Specific Heat

Specific heat describes the relationship of temperature to heat for a given substance. Water, with a specific heat of 4.184 J/g-degC has one of the highest specific heats of any substance. This property of water allows it to absorb or release large amounts of heat without changing its temperature dramatically. This is a very important feature, since it allows the large bodies of water on the earth's surface to moderate our climate so that temperatures do not get too hot or too cold. Water also has a very high heat of vaporization, and freezing. All of these factors combine to make the aquatic environment much more temperature-stable than our own terrestrial conditions. Marine organisms especially are adapted to fairly steady temperature conditions, and are often killed by abrupt changes in temperature that we would take in stride.

Seawater

Salinity

Seawater is a solution of salts of nearly constant composition, dissolved in variable amounts of water. There are over 70 elements dissolved in seawater but only 6 make up over 99% of all the dissolved salts; all occur as ions - electrically charged atoms or groups of atoms.

Major Constituents of Seawater

Chloride ( Cl- )

55.04% wt

Sodium ( Na+ )

30.61% wt

Sulphate ( SO4- )

7.68% wt

Magnesium ( Mg+ )

3.69% wt

Calcium ( Ca+ )

1.16% wt

Potassium ( K+ )

1.10% wt

Total

99.28% wt

Oceanographers use salinity - the amount ( in grams ) of total dissolved salts present in 1 kilogram of water - to express the salt content of seawater. Normal seawater has a salinity of 5 parts per thousand, also expressed as 35°. This equates to a specific gravity of 1.020 to 1.024.

Seawater varies in salinity from place to place, ranging between 34° and 37° in open ocean areas, down to 0° near the discharge of large rivers. High evaporation levels cause noticeably saltier surface water in the tropics, while freshwater runoff in some enclosed northern areas like the Baltic Sea dilutes the seawater to almost fresh. Seawater from Wormly in southern England is used as the international standard for seawater composition. Locally, the salinity of coastal waters averages about 31°, or a specific gravity of 1.022. The North Atlantic is the warmest and least saline of all oceans.

Daily average water salinity, as measured by Rutgers IMCS LEO-15 station ( off Tuckerton. )

Salinity measurements averaged by month shows a dip in winter, when freshwater
runoff from the land is high and temperatures and therefore evaporation are low.

As well as major elements, there are many trace elements in seawater - e.g., manganese (Mn), lead (Pb), gold (Au), iron (Fe), iodine (I). Most occur in parts per million (ppm) or parts per billion (ppb) concentrations. They are important to some biochemical reactions - both from positive and negative (toxicity) viewpoints.

Volcanic eruptions produce large volumes of gases that eventually reach the oceans -- most important are sulphate and chloride. Submarine eruptions at spreading ridges inject gases directly into the oceans; gases from sub-aerial volcanoes are dissolved in rainfall.

Chemical reactions between hot seawater and recently formed basaltic ocean crust lead to removal of magnesium and some sulphate from the seawater, while other elements like lithium and rubidium are added. Ocean water is circulated through the ocean crust by this exchange mechanism. The entire volume of all the ocean's water circulates every 5-10 million years. This is probably the main reason that the composition of seawater has been nearly constant over billions of years.

Many salts in seawater originate from weathering of rocks on land. As rocks are weathered to form soils, they release soluble constituents like silica and elements like sodium, calcium, potassium and magnesium. River waters also carry bicarbonate (HCO3-) - a by-product of weathering of silicate rocks or dissolution of limestone. Once they enter the oceans the dissolved salts remain, while the water continues to move through the hydrological cycle.

Because, there is geological evidence to suggest that seawater has retained the same salinity for billions of years ( at least 3.4 Ga ) other processes must be removing dissolved constituents from seawater. Some are removed by chemical reactions between seawater and the sediment ( e.g. manganese nodules. ) Others ( e.g. silica, nitrate, calcium carbonate ) are removed by organisms, including skeletal material. Other constituents, like sodium and chloride are removed as salts ( evaporite deposits ) when some of the seawater evaporates. This occurs when an arm of the sea is partially or totally cut-off from the open ocean and happens to lie in an arid climatic zone. The potash deposits of Saskatchewan were formed when an arm of the sea was isolated from the open ocean about 400 million years ago. Most of the water evaporated, leaving behind a thick deposit of salts. Wind blowing sea-spray also removes salt as aerosols which may be deposited along the coastline. Also, some minerals in seafloor sediments, such as clays, adsorb ( "grab" ) metallic ions from seawater onto their surfaces. When evaporites and seafloor sediments are uplifted, associated with collision of lithospheric plates, they again become subject to weathering and the dissolved salts eventually return to the sea.

pH

pH, or "percent Hydrogen", is a measure of the acidity or alkalinity of a substance, generally a water-based mixture. pH is measured on a scale from 0 to 14. The pH scale is logarithmic - the numbers are actually exponents, but you don't need to understand how that works to understand the concept.

A pH of 7.0 is by definition neutral. Acidity ( pH < 7.0 ) is a measure of the concentration of free hydrogen ions ( H+ or simply bare protons ) in a solution. Alkalinity ( pH > 7.0 ) is a measure of the concentration of "holes" where such a hydrogen ion could fit. Since water dissociates ( breaks apart ) into equal parts of hydrogen ions and "holes" ( in this case, hydroxide ions OH-, although any negative ion will act as a base ) it is equally an acid and a base, or perfectly neutral:

H2O <=> OH- + H+

Warning:Hydrogen hydroxide ( HOH ) is commonly used as an industrial solvent, and also in the manufacture of chemical and biological weapons. It is often present in cancerous tumors. Exposure to the gaseous form results in blisters and burns; contact with the solid form results in numbness followed by tissue damage; immersion in the liquid form has been linked to many fatalities.

In fact, another name for water is hydrogen hydroxide - HOH. ( Also: hydric acid, and that even more feared enviro-toxin: dihydrogen monoxide. )

Carbon dioxide is more soluble in water than most gases, owing to the polar nature of the CO2
molecule, similar in this regard to the water molecule itself. In combination with water, carbon dioxide readily forms carbonic acid, H2CO3. Carbonic acid may then dissociate to form bicarbonate ions HCO3- and hydrogen ions H+, which may further dissociate to form carbonate ions CO3-- and yet more hydrogen ions:

CO2 + H2O <=> H2CO3

H2CO3 <=> HCO3- + H+

HCO3- <=> CO3-- + H+

or

CO2 + H2O <=> H2CO3 <=> HCO3-
+ H+ <=> CO3-- + 2H+

Carbonate and bicarbonate are both well-known alkalines, or bases, while hydrogen ions are acidic. This series of reactions, which go equally well in either direction, forms the basis of a chemical buffering system which holds the pH ( or acidity ) of seawater to a stable and narrow range between 7.5 and 8.4. This is slightly alkaline, owing to the concentration of the strongly alkaline positive ions listed above: mainly calcium carbonates, but also similar compounds of potassium and sodium.

By comparison, pure distilled water is neutral in pH ( by definition, pH = 7.0 ) whereas most natural freshwater bodies tend toward acidity, in the range from 5.0 to 7.5. This acidity is due mainly to products of biological decomposition. In some areas, compounds leaching from fallen tree leaves, especially oak, can make the water extremely acidic, with a tea-like coloration. This is common in many areas of southern New Jersey.

For comparison, the pH of lemon juice is about 3.0; the pH of soap is about 10.

The Mid-Atlantic Bight and the Gulf Stream

The Mid-Atlantic Bight is the region enclosed by the coastline from the North Carolina Capes in the south to Cape Cod in the north. The outer boundary is typically taken as the edge of the Gulf Stream. These latitudes are classified as cold temperate, with wide seasonal variations in temperature and solar radiation.

The flow of the Gulf Stream in the North Atlantic

Of note is that the current diverges from the continental margin off the Carolinas, and passes far offshore of New Jersey. The Gulf Stream is primarily a surface current, which floats atop the colder denser deep ocean waters.

Schematic representation of the Gulf Stream

This diagram shows how loops are pinched off, forming circulating pockets of tropical water which spin away from the main flow on either side. Some of these bring warm tropical water to our shores, even though the main flow is hundreds of miles out to sea and heading for England. ( Cyclonic simply means counterclockwise, the normal direction of rotation of atmospheric storms in the northern hemisphere. )

Temperature

A satellite sensing image of sea surface temperatures

This image shows the warm Gulf Stream waters ( red ) meandering offshore, leaving the Mid-Atlantic Bight area filled with the cold water of the northern Labrador Current. Note the eddy of warm water ( yellow ) swirling away from the main flow at middle right. The water temperatures shown here are unusually cold for this time of year.

A closer-in view of the area that is of interest to divers.

Note the bottom contours, and the finger of warm water at the lower right. These warm surface waters float above colder layers below. Water composition is also influenced by freshwater river flows and runoff from the land, as is evidenced by the temperature gradients along the shore and in the shallow bays.

Monthly average surface water temperature data from the Long Island weather buoy

Monthly average coastal bottom temperature

Note how bottom temperature trends along with offshore surface temperature ( above ) peaking somewhat later ( September vs. August. ) the LEO-15 stations are inshore in relatively shallow water, and these temperatures should not be considered indicative of the entire area. In particular, deep waters offshore are considerably colder in the summer months.

Daily average bottom water temperatures over several years,
as measured by Rutgers IMCS LEO-15 station ( off Tuckerton. )

On a related subject, here is a plot of water temperatures at a popular
freshwater dive site -
Dutch Springs quarry. Note that the data covers
only part of the year, from April to November.

Thermoclines & Stratification

In temperate-zone seas as off the New Jersey coast, the amount of sunlight varies seasonally. As a result, the amount of solar energy entering the water varies, which in turn alters the temperature in the upper water layers. The thermal structure of the water column thus changes seasonally.

Seasonal thermocline structure in temperate seas

In the summer months the sun is high, days are long, and the upper layers heat up and become less dense than underlying layers. In other words, the water column becomes thermally stratified and no mixing occurs. In the fall the amount of solar energy entering the water column decreases, days become shorter, upper layers cool, and thermal stratification decreases. Finally, a point is reached where the temperature of the surface layers has been reduced to such an extent that the density of the layer is little different from that of the underlying mass. At this point, mixing can occur whenever sufficient wind is available.

In winter, usually the storm season in the temperate zone, the sun is lowest on the horizon, solar energy input to the water is at a minimum, thermal stratification is at a minimum or absent, and mixing occurs. With the onset of spring, the days become longer, the solar energy increases, the upper layers begin to rise in temperature, and the system moves toward re-establishment of thermal stratification.

This same sequence of events also occurs in many freshwater bodies of water.

Oxygen & Other Dissolved Gases

Seawater also contains small amounts of dissolved gases ( nitrogen, oxygen, carbon dioxide, hydrogen, and trace gases. ) Water at a given temperature and salinity is saturated with gas when the amount of gas entering the water equals the amount leaving during the same time. Surface seawater is normally saturated with atmospheric gases such as oxygen and nitrogen. The amount of gas that can dissolve in seawater is determined predominantly by the water's temperature and salinity. Increasing the temperature or salinity reduces the amount of gas that can be dissolved.

35°F seawater at the surface typically contains about 0.54% dissolved oxygen by volume at 68°F, and about 0.80% at 32°F ( compared to 21% for air ! ) Most fish become stressed when dissolved oxygen levels fall to 0.2-0.4%, and large-scale die-offs of aquatic fauna may occur at levels below 0.2%.

Dissolved Oxygen

Temperature

Freshwater

Saltwater 35°

deg C

deg F

saturated
(100%)

minimum
healthy
( 6 mg/l )

saturated
(100%)

minimum
healthy
( 5.5 mg/l )

0

32

14.6 mg/l

41%

11.7 mg/l

47%

5

41

12.8 mg/l

47%

10.4 mg/l

52%

10

50

11.3 mg/l

53%

9.3 mg/l

58%

15

59

10.1 mg/l

59%

8.5 mg/l

65%

20

68

9.1 mg/l

66%

7.8 mg/l

71%

25

77

8.2 mg/l

73%

7.1 mg/l

77%

30

86

7.5 mg/l

80%

6.5 mg/l

85%

A rule of thumb is that for any cold-blooded animal, the metabolic rate roughly doubles for every 10°C ( 18°F. ) Thus, a cold-blooded organism will require twice as much oxygen at 77°F as it would at 59°F, while at the same time, the available oxygen has dropped by approximately 20% ( likely more. ) Many aquatic creatures ( fishes especially ) are capable of greatly varying their respiration rate to adjust to a range of temperatures and oxygen levels.

The amount of dissolved oxygen in the water is one of the principal indicators of the health of an aquatic ecosystem. If dissolved oxygen levels drop too low, known as a Low Oxygen Event, animals which can move ( e.g., fish and lobsters ) will leave, and animals which can't ( e.g., clams and oysters ) will become stressed, and eventually die if the levels do not increase, an event known as a fishkill. There are some environments which have naturally occurring cycles of low oxygen, and some animals have evolved to tolerate the low oxygen. In some areas where the nearby coastline is heavily populated, the severity and frequency of low oxygen conditions can be directly attributed to human activities.

The naturally occurring processes which affect the oxygen cycle are:

The sun and nutrients stimulate photosynthesis in marine plankton and a bloom occurs.

As the plankton die, they rain down to the bottom and begin to decompose. The microorganisms which decompose the dead organic matter start to consume the available oxygen.

Mixing of new dissolved oxygen from the upper waters into the lower water column is usually inhibited because the sun has also warmed the upper layer, creating a thermocline which inhibits mixing. The microorganisms continue to use up the oxygen.

As the dissolved oxygen concentration drops below certain levels, fish and lobsters start to move out. Shellfish are unable to escape the low oxygen levels, they will die unless the oxygen levels go up again due to a major mixing event ( storms ) or the bloom ends and the number of microorganisms using up the oxygen decreases.

Low oxygen events like this are usually short-lived and the marine community recovers quickly. However, human processes can aggravate matters:

The runoff from land now carries extra nutrients from lawn fertilizers and domestic animals, and effluent pipes from treatment plants add sewage and additional detrital material.

The plankton bloom can grow faster and last much longer because of the extra nutrients.

More dead phytoplankton and detritus is providing more material for the microorganisms on the bottom to decompose, intensifying and prolonging the dissolved oxygen depletion.

Once water sinks below the ocean surface, dissolved gases can no longer exchange with the atmosphere. The amount of gas in a given volume of water may remain unchanged, except by movement of gas molecules through the water - diffusion ( slow process ) or by the water mixing with other water masses containing different amounts of dissolved gas. In general, nitrogen and rare inert gases ( argon, helium, etc. ) behave this way - their concentrations are conservative and only affected by physical processes. In contrast, some dissolved gases are non-conservative
and actively participate in chemical and biological processes that change their concentrations. Examples are oxygen and carbon dioxide - released and used at various rates in the oceans, especially by organisms.

Nitrogen, essential for plant growth, is even less soluble in water - only 0.00005% in the richest waters - 1/10,000 of the level typically found in terrestrial soil.

Plankton & Visibility

Thermal stratification has a pronounced effect on the growth of phytoplankton. Phytoplankton are single-celled algae suspended in the water column. Phytoplankton require chemical nutrients, primarily invisible phosphates and nitrates, dissolved at fairly low levels in the water, as well as oxygen. Phytoplankton also requires sufficient light for photosynthesis in order to grow. This restricts its growing zone to the upper layers of the water column, known as the photic zone. While it does not immediately die in the darker deep waters, phytoplankton cannot grow there. Phytoplankton are fed upon by zooplankton ( microscopic animals ) thus forming the first two levels in the food pyramid of the sea. Although almost all of these organisms are too small to see with the naked eye, all together they can have a great effect on water clarity, or what we divers call visibility.

Nutrient dispersal in the water column

In the winter, surface and deep waters are well mixed, such that sufficient nutrients are found at all levels, including the surface waters that receive enough light for phytoplankton to grow. The amount of plankton that grows in the winter is limited by the cold temperatures and the relatively poor winter illumination. As the days lengthen into spring, phytoplankton make use of the increased light and high nutrient levels, resulting in an "algae bloom." the zooplankton lag behind the phytoplankton in numbers, but eventually they catch up and the "bloom" is consumed.

With the formation of the summer thermocline, the surface waters are separated from the bottom waters, making the nutrients there unavailable to the phytoplankton, which soon exhaust what nutrients are available to them in the lighted upper waters. At this point, numbers of both phytoplankton and zooplankton drop. For us, the water clears up. This situation is maintained until the thermocline breaks down in the fall, allowing the untapped nutrients from the deep waters to mix up into the photic zone again, and generally resulting in another algae bloom, although not as great as in the spring. As with stratification, these same algae cycles also occur in many freshwater bodies as well.

Distribution of chlorophyll: greener = more chlorophyll = more phytoplankton

While almost all phytoplankton is microscopic, not all zooplankton is so small - some jellyfishes are up to 8 feet across ! Many planktonic crustaceans, worms, fish larvae, and other creatures are visible to the naked eye, if you are looking.

Marine Snow

Another factor that affects water clarity and visibility is "marine snow." Marine snow is composed of microscopic bits of non-living matter - soot, dust, sand, dead plankton, fecal particles, etc - that is held together by sticky bacterial waste products and slowly sinks down in the water column.

Marine snow is extremely fragile, and at night the swimming actions of feeding zooplankton breaks it up. However, during the day it may aggregate into visible clumps or strings, what we divers refer to as "Egg Drop Soup." Marine snow was actually discovered by scuba divers, since it is so difficult to collect by the usual means of plankton sampling because it is so fragile.

Sea Foam

When it looks as if the beach or surf is littered with billows of soapsuds, the reason is not that someone has dumped a carload of detergent into the ocean. This common sight of spring and summer ( and occasionally, brief periods in fall and winter ) results from prolific the reproduction of "phytoplankton." Each microscopic individual is housed in a skeleton made of calcium or silica. Conditions of sun, temperature, and nutrients stimulate rapid growth and reproduction. When the cycle has run its course, billions of individuals die.

Sea foam is created when waves or strong winds inject air into the dissolved organic matter in ocean water, forming bubbles. The organic matter, mostly made of dead phytoplankton, contains protein that gives the salt water enough surface tension to form bubbles. Surf and winds cause the mass to pile up in the familiar, suds-like masses. Windrows of phytoplankton remains show that the sea off our coast is producing tons of food for other creatures in the food web.

Upwellings & Fishkills

In coastal areas, strong offshore ( land to sea ) winds can push surface waters away from the land. Cold nutrient-rich water then flows up from the deep up to replace it, thermocline or not, in an event known as an upwelling. Upwellings usually result in localized algae blooms, which are sometimes intense enough to cause low-oxygen events and even fishkills. Such a fishkill, of unprecedented proportions, occurred in 1976, affecting almost the entire marine environment off the New Jersey coast.

During the summer months, the surface of the ocean near our coast is heated by the sun. This warming causes stratification (warm surface/cold bottom). Typically, winds during our summer months are from the southwest, bringing all that hot humid air up from the Gulf of Mexico. These winds do not blow that surface layer directly to the northeast, but to the southeast. This 90 degree difference in wind and water current direction is due to the spin put on the water by the earth's coriolis force.

When the warm surface water is blown offshore, the cold bottom water rises, like a conveyor belt, and hits the beach. This cold water also brings sediment up from the bottom. Phytoplankton (microscopic plants) which float along the surface, used this sediment as food and bloom, causing the water to become green and murky. The figure above demonstrates what happens at the surface and below the water when winds blow from the southwest along the New Jersey coast. This is why the beaches can be so cold and "dirty" on the warmest days of the year.

compiled from various sources

The Ocean Fishkill of 1976

by Bill Figley, Jeff Carlson, Dan Vaughan, Sue Hollings

If you are an avid marine fisherman, you probably will not forget the summer of 1976 - dead and dying fish in the surf and inlets; wrecks devoid of fish, crabs, lobsters, and other marine organisms; disrupted fish migrations; brown scum washing up on beaches; and extensive streaks of red water along the surf.

The first indications of trouble were received from sport divers who found dead and dying fish and shellfish on wrecks located off the northern coast of New Jersey in late June of 1976. The divers noted that the bottom waters were extremely murky and that a dark scum covered the wrecks and bottom. In response, the National Marine Fisheries Services ( NMFS ) from the Sandy Hook Lab and the Marine Fisheries Section of DEP's Division of Fish, Game and Shellfisheries embarked on sampling programs designed to assess water quality in the reported area. Water samples indicated severely low concentrations of dissolved oxygen in bottom layers. The murky water and bottom scum were found to be the result of a thick bloom of algae that consisted primarily of Ceratium tripos. Soon, reports of similar mortalities were being received from many offshore areas as far south as Atlantic City.

More than 3,000 square miles of New Jersey's coastal waters suffered severe depletion of dissolved oxygen near the bottom. During July and August of 1976, bottom-dwelling fish were absent from an area comprising more than 1,100 square miles.

As the seriousness of the problem became more evident, sampling efforts by the NMFS and the Division were expanded to cover a greater portion of the New Jersey coast. These investigations indicated that bottom waters in an area of 3,000 square miles extending from four to about sixty miles offshore and from Sandy Hook to below Atlantic City were either devoid of oxygen ( anoxic ) or had very low ( hypoxic ) concentrations of dissolved oxygen. Slow-moving and attached organisms, such as crabs, clams, scallops, starfish, worms, and barnacles, perished in areas where the anoxic conditions persisted. Fortunately, most fish were able to sense the gradually declining oxygen levels and avoided pockets of anoxic water. However, fish species such as eel pout, cunner, and sea bass, which depend upon wrecks and other obstructions for protection, were reluctant to leave such shelters and consequently suffered high mortalities. Through sample trawling, NMFS biologists found that fish were completely absent from more than 1,100 square miles of ocean bottom during August.

The impact of the "bad water" on the State's marine fisheries resources was disastrous. About 69 percent ( 19 million bushels ) of the State's most valuable fishery stock, the surf clam, was destroyed; the potential economic value of this lost resource amounted to more than $430 million. Losses to lobster industry amounted to more than $1 million. The inshore trawler fishery lost more than $5 million. Bluefish avoided the bad water and did not venture into coastal waters off northern New Jersey. Tuna remained well offshore. Bottom species - fluke, sea bass, and ling - had to vacate traditional offshore grounds. Charter and party boats dependent on these species lost future fares from hesitant and disappointed fishermen, had to cancel trips, and were forced to use more fuel to reach the dispersed fish. In all, the State's charter and party fleet lost an estimated $1.7 million during 1976. In addition, biologists suspect that future stocks may have been damaged because of the disruption of the spawning activities of many species and of the mortality of fish eggs and larvae drifting through anoxic waters.

An Explanation of the Kill

Division biologist collecting a
water sample off Barnegat Light.

The immediate causes of fish and shellfish mortalities were the lack of sufficient concentrations of oxygen in bottom waters and the presence of poisonous hydrogen sulfide gas in anoxic areas. The depletion of dissolved oxygen occurred in the following manner:

During the late spring, ocean water temperatures vary little from surface to bottom. Oxygen reaches surface waters through diffusion from the air and is also produced by algae during photosynthesis. It is circulated through the water column by diffusion, wave action, and various currents. As the sun warms the surface layers and the colder, heavier waters sink to the bottom, two layers, one of warm water and one of cold water, are formed.

Between the layers is a thin zone called the thermocline in which temperature declines rapidly with depth. This condition is especially noticeable while swimming in lakes and ponds - diving underwater, the swimmer will suddenly encounter an abrupt drop in temperature. During 1976, the thermocline extended from 30-to 60-foot depths. Since the bottom layer is significantly heavier than the surface, the two layers resist mixing ( like oil and vinegar. ) the thermocline acts as a barrier to diffusion of oxygen between the surface and bottom. Thus the oxygen supply trapped in the bottom layer must be sufficient to sustain all of the living organisms through the summer, and normally it is capable of doing this. During 1976, however, the decay of tremendous quantities of Ceratium algae which died and slowly sank to the bottom, and the chemical and bacterial decomposition of other organic matter, placed additional demands on the limited oxygen supply. During July and August of 1976, most of the bottom waters off the coast of New Jersey had dissolved oxygen levels below 1.4 milliliters per liter ( mI/I ), the concentration below which most fish species begin exhibiting signs of serious stress, and a large area situated off Long Beach Island had absolutely no dissolved oxygen.

Hypothetical cross-section of ocean depicting the changes in water temperatures and dissolved oxygen concentrations which occur after the formation of a thermocline and the unusual hypoxic condition which led to the fishkill of 1976.

Inshore waters less than 60 feet deep were generally not affected because no thermocline was present to trap bottom waters and wave action was able to circulate oxygen throughout the water column. Some species were forced to spend the summer within this narrow band along the shore. At many times during the summer, sport fishermen were able to make large catches ( perhaps too large ) of fluke in the inlets or just outside the surf.

What caused the anoxic water condition?

Ceratium tripos is a dinoflagellate alga
which experienced a tremendous
population bloom during 1976.

The immediate response of sport and commercial fishermen was that the kill was caused by sludge dumping. They believed that the brown and black scum, which covered wrecks and occasionally washed up onto beaches, was raw sewage. In most cases, however, the scum was found to be gelatinous masses of decayed algae.

NMFS biologists believed the anoxic water condition was due in large part to unusual weather conditions. First, surface air temperatures were unusually high during the first six months of 1976 ( 2-5 F above normal, ) resulting in a more rapid warming of surface waters. Second, river runoff began two months earlier in the spring than usual. Third, during February and March, the predominant winds were from the south. Finally, storm activity was very low during the spring and summer. All these factors favored the early formation of a thermocline. Consequently, the oxygen supply below the thermocline was subjected to two additional months of oxygen-using activity by marine organisms. By late May, oxygen concentrations were as low as they normally are in July. The combination of early thermocline formation and the decomposition of tremendous quantities of Ceratium algae and other organic matter resulted in the eventual exhaustion of oxygen supplies in many areas off the New Jersey coast.

With regard to the fishermen's theory, no one is really certain what role sludge dumping played in the anoxic water condition of 1976. The ocean off New York and New Jersey now receives nutrients ( phosphates and nitrates ) and organic material from many sources - the dumping of sewage sludge, dredge spoil, and chemical wastes; municipal sewage outfall; river runoff which includes treated and untreated sewage; industrial wastes and farm fertilizers; sewage from pleasure and commercial boats; and more. The nutrients act as fertilizers and may stimulate algae blooms; then, as the algae die and sink to the bottom, their bodies are decomposed by bacteria. Other organic materials in the water are decomposed chemically and by bacteria. The decomposition process utilizes the oxygen supply that is necessary for all marine organisms commonly associated with clean water.

The Division's Program

The future of New Jersey's economically valuable commercial and recreational marine fisheries and possibly its ocean-based tourist trade is dependent upon the health of the State's coastal waters. At present, no one is certain of the ultimate cause of the anoxic waters of 1976 or if this condition will reoccur in future years.

In an attempt to gather the data needed to identify more specifically the primary factors responsible for this problem, DEP's Division of Fish, Game, and Shellfisheries continued where it left off in 1976 with an intensive ocean sampling program through the summer and fall of 1977. The primary objectives of the sampling cruises were to monitor dissolved oxygen concentrations, temperatures, and nutrient levels of nearshore ocean waters. Fortunately, conditions were improved over those of 1976. No areas were sampled that were completely devoid of oxygen. However, a large area extending from Manasquan to Beach Haven between one-half and twelve miles offshore had dissolved oxygen concentrations between 1 and 2 mI/I during August. These levels are considered critical, but not lethal, to fish and invertebrate life.

Another phase of the study was to investigate the flow of nutrients and organic material from northern inlets into the ocean. Of the four inlets sampled ( Barnegat, Manasquan, Navesink and Raritan River, ) the Raritan River had much greater inputs of nitrogen and carbon compounds.

During the spring of 1978, sampling operations will continue with special emphasis on monitoring algae and benthic ( bottom-dwelling organisms ) populations. The benthic survey will provide the data needed to assess the recovery of the resident populations that were decimated by the anoxic water.

Determining the primary factors that led to the anoxic water condition off our coast is only the initial step in safeguarding one of New Jersey's most valuable natural resources.

This article first appeared in New Jersey Outdoors - July / August 1978

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